Question

Determine which of the molecules are polar. Explain your answers. a. \(C-O\) b. Si- O c. Ge \(-0\) d. \(C-C l\) e. \(C-B r\)

Short answer

The polar molecules among the given options are: \(C - O\), \(Si - O\), \(Ge - O\), \(C - Cl\), and \(C - Br\). All of these molecules have an electronegativity difference greater than 0.4, making them polar. Additionally, since the molecules are diatomic, there is no molecular symmetry to cancel out the polarity.

Step-by-step solution

Defining Polar Molecules

Polar molecules are those with an uneven distribution of electrons due to the electronegativity difference between the atoms. This results in a partially positive and partially negative charge at either end of the molecule.

Examine Electronegativity Differences

Look up the electronegativity values for Carbon (C), Oxygen (O), Silicon (Si), Germanium (Ge), Chlorine (Cl), and Bromine (Br). Calculate the difference in electronegativity for each molecule: a. C-O b. Si-O c. Ge-O d. C-Cl e. C-Br

Identifying Polar Bonds

If the electronegativity difference is greater than 0.4, the bond is considered polar. Compare the differences calculated in Step 2 to this threshold to determine the polarity of the bonds.

Analyze Molecular Geometry

If the molecules are symmetrical, their polarity may cancel out. In this case, all of the molecules are diatomic, and hence there is no molecular symmetry to evaluate.

Determine the Polarity of the Molecules

a. \(C - O\) has an electronegativity difference greater than 0.4, so it is a polar molecule. b. \(Si - O\) has an electronegativity difference greater than 0.4, so it is a polar molecule. c. \(Ge - O\) has an electronegativity difference greater than 0.4, so it is a polar molecule. d. \(C - Cl\) has an electronegativity difference greater than 0.4, so it is a polar molecule. e. \(C - Br\) has an electronegativity difference greater than 0.4, so it is a polar molecule. Based on the analysis, all of the given molecules should be considered polar due to significant electronegativity differences between the bonded atoms.

Key concepts

Electronegativity

Electronegativity is a measure of how strongly an atom can attract or pull on electrons when forming a bond. Electrons are the basis of chemical bonds, and their distribution within a molecule significantly affects the properties of the molecule, including its polarity. The electronegativity scale was developed by Linus Pauling, and it allows chemists to predict the types of bonds that will form between different elements.

Atoms with high electronegativity, such as Oxygen (O), Chlorine (Cl), and Bromine (Br), have a stronger pull on shared electrons. In contrast, atoms with lower electronegativity, like Silicon (Si) and Germanium (Ge), exert a weaker pull. When there is a significant difference in electronegativity between two bonded atoms, the electrons in the bond are not shared equally. This leads to the formation of a dipole, with one end of the bond having a partial negative charge and the other a partial positive charge.

Understanding Polarity Through Electronegativity

To determine the polarity of a bond, look at the electronegativity values of the bonded atoms. A difference greater than 0.4 usually suggests that the bond is polar. For example, a Carbon (C) atom bonded to an Oxygen (O) atom will form a polar bond, as Oxygen is more electronegative and will attract the electrons more strongly than Carbon.

Chemical Bonding

Chemical bonding involves the joining of atoms to form new substances. The type of bond that forms between atoms depends on their electronegativity and valence electrons—those found in the outermost shell of an atom. There are three main types of chemical bonds: ionic, covalent, and metallic.

Ionic bonds occur between metals and nonmetals when one atom donates electrons to another, resulting in positively and negatively charged ions. Covalent bonds form when atoms share electrons to reach a stable electron configuration, and they can be nonpolar or polar. Nonpolar covalent bonds occur between atoms with similar electronegativities, where the electron sharing is quite equal. However, polar covalent bonds are the result of atoms with different electronegativities sharing electrons unequally, as seen in the bond between Carbon (C) and Oxygen (O) in the CO molecule.

Role of Chemical Bonding in Molecular Polarity

Each bond's polarity contributes to the overall polarity of the molecule. However, the molecular geometry also plays a crucial role. In diatomic molecules, such as those in the exercise, the polarity is solely determined by the bond as there is no geometric symmetry to consider. As such, all the bonds listed (C-O, Si-O, Ge-O, C-Cl, C-Br) are polar due to their chemical bonding nature—polar covalent.

Molecular Geometry

Molecular geometry pertains to the three-dimensional arrangement of atoms within a molecule. The spatial configuration has a profound effect on the physical and chemical properties of the substance, including the molecule's polarity. When considering molecular geometry, it's essential to look at the molecule's symmetry and the arrangement of polar bonds.

In symmetrical molecules, like methane (CH4) or carbon dioxide (CO2), the individual dipole moments can cancel each other out, leading to a nonpolar molecule. On the other hand, asymmetrical molecules will likely be polar because the dipole moments do not cancel out. Molecules with bent or trigonal pyramidal geometries, such as water (H2O), are classic examples of polar molecules.

Molecular Geometry and Dipole Moments

For exercises involving diatomic molecules, polarity is straightforward, as there is no complex geometry to consider. A single polar bond makes the entire molecule polar. In more complex molecules, even if there are polar bonds present, the overall molecular polarity must be considered through the lens of molecular geometry. Predicting these requires a thorough understanding of VSEPR (Valence Shell Electron Pair Repulsion) theory, which can explain the molecular shapes based on the repulsion between electron pairs.